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Oxides


All the elements in Group 2 burn in O2 to form metallic oxide. BeO is usually made by burning of the metal, but the other metal oxides are usually obtained by thermal decomposition of the carbonates, MCO3.

The alkaline earth metals are less electropositive than the alkali metal, thus the oxides of alkaline earth metals are less basic than those of alkali metals. On descending the Group 2, the oxides become more basic due to increase in the electropositive character of the metals.

BeO is an amphoteric oxide. It dissolves in acids to give salts, and in alkalis to give beryllates, which on standing precipitate as hydroxides. Peroxides are formed with increasing ease and increasing stability as the size of metal ions becomes larger.

BaO2 and SrO2 can be made by passing air over BaO at elevated temperature and pressure.
MgO2 and CaO2 is made by reacting Ca(OH)2 with H2O2 and then dehydrating the product. No peroxide of beryllium is known.

Hydroxides


Except Be(OH)2, all other hydroxides are basic. The base strength increases on descending the group.
Be(OH)2 is an amphoteric hydroxide.
Mg(OH)2 is a weak base.
Ca(OH)2 and Sr(OH)2 are moderately strong bases.
Ba(OH)2 is a very strong base.
Solution of Ca(OH)2 is known as lime water while that of Ba(OH)2 is known as baryta solution. Both these solutions are used to detect carbon dioxide.
Ca(OH)2 + CO2 CaCO3 + H2O    Ca(HCO3)2
                           lime water                   milkyness disappear
                           turns milky

 

The solubility of hydroxides increases on descending the group (fig. 10.15).


Explanation
The solubility of an ionic salt depends on the following two factors.

  1. Lattice Energy
    It is the energy released when the requisite number of gaseous ions are converted into 1 mol of crystalline lattice. Higher the lattice energy more tightly the ions are held together in the crystal, less will be its tendency to split into ions to pass over to the solution, i.e. less will be its solubility.
  2. Hydration Energy
    It is the heat released when the dissolved ions get hydrated in the solution. Higher the hydration energy, the more heat released and thus higher is the solubility.
    Hence,
    Higher the lattice energy, lesser the solubility and higher the hydration energy, more the solubility.
In case of hydroxides, the size of anion is not large, the change in the size of cation causes significant change in lattice energy which decreases from Be2+ to Ba2+ ions. This factor is predominant as compared to the decrease in the hydration energy, thus causing the solubility of hydroxides to increase as we go from Be to Ba salts.

Carbonates and Bicarbonates


Alkaline earth metals form solid carbonates. Bicarbonates are known only in solution because of the less basic character of alkaline earth metals. The thermal stability of carbonates increases on descending the group. The decomposition temperatures are as follows:
    BeCO3        MgCO3        CaCO3          SrCO3       BaCO3
< 100 °C        540 °C        900 °C         1290 °C    1360 °C

Explanation
As the size of M2+ ion is increased on descending the group, its polarization power on CO32- decreases making carbonate more stable. The solubility of carbonates decreases on descending the group. In case of carbonates, the size of CO ion is large. The changes in the size of metallic ions from Be2+ to Ba2+ do not make much difference in the lattice energy. However, the change in hydration energy is more predominant which decreases regularly from Be2+ to Ba2+ causing the decrease in the solubility as we go from Be to Ba.

Sulphates


The solubility of the sulphates of Group 2 metals decreases down the group.
BeSO4 and MgSO4 are soluble.
CaSO4 is sparingly soluble.
SrSO4, BaSO4 and RaSO4 are virtually insoluble.

Explanation
In case of sulphates, the size of SO42- is large. The changes in size of metallic ions from Be2+ to Ba2+ do not make much difference in the lattice energy. However, the changes in hydration energy is more predominant which decreases regularly from Be2+ to Ba2+ causing the decrease in the solubility as we go from Be to Ba.

All sulphates decompose on heating;
MSO4 MO + SO3

The thermal stability increases on descending the group. The decomposition temperatures are
BeSO4              MgSO4             CaSO4         SrSO4
500 °C             895 °C             1149 °C        1374 °C

Explanation
As the size of M2+ ion is increased on descending the group, its polarization effect on SO42- decreases making sulphate more stable.

Nitrates


Alkaline earth metals form hydrated nitrate when crystallized from aqueous solutions. Heating the hydrated solids does not give the anhydrous nitrate because the solid decomposes to the oxide.

Anhydrous nitrates can be prepared by treating chloride with liquid N2O4 in ethyl acetate. Beryllium is an exception which gives a basic nitrate in addition to the anhydrous nitrate.

Becl2 Be(NO3)2.2N2O4 Be(NO3)2 [Be4O(NO3)6] N2O4

       
 Basic beryllium acetate, [Be4O(CH3COO)6], has a similar structure.

Halides


Alkaline earth metals combine directly with halogens to form halides of the type MX2 at the appropriate temperature. They can also be prepared by treating metals or their carbonates with halogen acids.

Beryllium chloride is readily prepared from the metal oxide
BeO + C + Cl2 BeCl2 + CO

It cannot be prepared in aqueous solution due to the formation of hydrated ion [Be(H2O)4]2+. The halides obtained has the molecular formula [Be(H2O)4]Cl2, [Be(H2O)4]F2, etc. It cannot be dehydrated by heating as it undergoes hydrolysis.
[Be(H2O)4]Cl2 Be(OH)2 + 2HCl + 2H2O

Beryllium halides are covalent and fume in air due to hydrolysis:
BeCl2 + 2H2O Be(OH)2 + 2HCl

In vapour phase it is present as BeCl2 and (BeCl2)2 and in solid phase it is polymerized.
Cl—Be—Cl                            
monomer                                      dimer

In the polymerized beryllium chloride, a halogen atom bonded to one beryllium atom uses a lone pair of electrons to form a coordinate bond to another beryllium atom. Two such halogens are present in between two beryllium atoms.

The halides of calcium, strontium and barium are essentially electrovalent and dissolve in water to give neutral solutions. They are hygroscopic and form hydrates like CaCl2 6H2O and MgCl2 6H2O. Anhydrous CaCl2 is a well-known drying agent. Anhydrous MgCl2 is used in the electrolytic extraction of magnesium.

Complexes


Alkaline earth metals have a stronger tendency to form complexes than do alkali metals. This is because of their smaller size and high charge. The tendency of complex formation decreases down the group.

Beryllium forms many complexes like BeF42-, [Be(H2O)4]2+, [Be(C2O4)2]2-, [Be4O(Ac)6], etc. Of the others, only Mg and Ca show more tendency to form complexes in solution, and these are usually with oxygen-donor ligands.

In most cases beryllium is tetra coordinated in complexes and the tetrahedral arrangement adopted corresponds to sp3 hybridization of beryllium orbitals. These are the orbitals (2s and three 2p) which are available in the valence shell of beryllium atom. The structures of beryllium oxalate complex and basic beryllium acetate are as follows:

The most important complex formed by magnesium is chlorophyll, the green plant pigment. It is vital for the photosynthesis of sugars in plants. Both magnesium and calcium form complexes with strong complexing agents like ethylenediamine tetraacetate (EDTA).




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